Periodic Trends in Ionic Radii in Modern Periodic Table

Q: What is the magnitude of ionic radius of of magnesium ?

The ionic radius of magnesium ion (Mg 2+ ) is 173 pm.

Q: What is ionic radius of Chlorine ion ?

The ionic radius of Chlorine anion (Cl − ) is 175 pm.

Periodic Trends in Ionic Radii in Modern Periodic Table - Atomic Radius with FAQs

Edited By Team Careers360 | Updated on Jul 02, 2025 04:52 PM IST

What is a Periodic Table?

A periodic table is not anything, however the association of all of the recognised factors in a way in which the factors having comparable properties are categorised collectively in tabular form.

The improvement of the periodic desk passed off withinside the following series:-

Lavoisier Classification-

This Story also Contains

What is a Periodic Table?
Lavoisier Classification-
2. Prout’s Hypothesis-
3. Dobereiner Triad Rule-
4. Newland Octave Rule-
5. Lother Meyer's Curve-
6. Mendeleev's Periodic Table-
7. Modern Periodic Table -
Atomic Radius –
Ionic Radius trend
Cationic Radius-
Anionic Radius-
Ionic size or Size of isoelectronic species:-
Periodic trends in ionic radius-
Periodic trends in ionic radius across the period:-
Factors Affecting Atomic Size:-
In a Period-
In a Group-
Lanthanide Contraction-
S-Block variation-

The category turned into truly primarily based totally on factors being metals and non-metals.

Metals being the factors tending to lose electrons and non-metals being the factors tending to just accept electrons.

Also read -

2. Prout’s Hypothesis-

This speculation turned into primarily based totally on the presumption of all of the factors being made from hydrogen.

3. Dobereiner Triad Rule-

Three detail corporations have been made with the aid of using him having the equal chemical properties; in which the atomic weight of a center element is almost identical to the common weight of the first and third elements

4. Newland Octave Rule-

The factors have been organized in increasing order in their atomic loads and look at that the properties of each eighth detail have been just like the first detail.

5. Lother Meyer's Curve-

He plotted a curve among atomic extent and atomic weight of various factors and concluded that the physical properties of elements are periodic capabilities in their atomic weight.

6. Mendeleev's Periodic Table-

Mendeleev's periodic regulation states that the bodily and chemical houses of the factors are the periodic feature in their atomic weight. He turned into the primary scientist to set up the factors in a scientific way i.e. in horizontal and vertical columns.

7. Modern Periodic Table -

The Modern Periodic table turned into proposed with the aid of using Moseley primarily based totally on atomic number.

Moseley concluded that the physical and chemical houses of factors are periodic function in their atomic number. This manner that once the factors are organized in growing order in their atomic number factors having comparable properties receives repeated after normal intervals.

Atomic Radius –

The common distance of valence shell electrons from the nucleus is known as the atomic radius.

It may be very tough to degree atomic radius due to the fact the isolation of a single atom is tough. Also, there's no well-described boundary for the atom.

So, the atomic radius may be appropriately described as-

1. Half the internuclear distance (d) among atoms in a homoatomic molecule.

2. This internuclear distance is likewise known as bond length.

Based at the chemical bonds, the atomic radius is split into 4 categories:-

a. Covalent radius

b. Ionic radius

c. Metallic radius

d. Van der Waal’s radius

Let’s see what is ionic radius definition is-

Ionic radius definition

The radius of an atom's ion is called the ionic radius. Ionic radius is measured by dividing the lengths between the two nuclei of the respective ions based on their sizes. Hence, ionic radii are nothing but the distance from the nucleus of an ion up to which it has an influence on its electron cloud.

Ionic Radius trend

Ionic radius increases on moving top to bottom in periodic table.

Ionic radius is further classified into cationic and anionic radius. Table of ions and ionic radius table help us to know better examples of ions and its radius.

Cationic Radius-

A cation is a positively charged ion formed on losing electron(s) in a neutral state.

Cationic radius is always smaller than atomic radius because after losing electron(s), the number of electrons reduces, by the number of origins remains the same. This happens due to the increase in effective nuclear charge (Zeff). Hence, electrons are pulled towards the nucleus and the atomic radius decreases moreover after losing all the electrons from the outermost shell, the penultimate shell becomes the ultimate shell which is nearer to the nucleus so size decreases.

Size of cation ∝ 1/the magnitude of charge or Zeff

Eg : Fe > Fe+2 > Fe+3

Pb+2 > Pb+3

Anionic Radius-

Anionic radius is a negatively charged ion formed on gaining electron in the neutral state.

Anionic radius is always larger atomic radius because in an anion, the number of electrons are more than the umber of protons and this leads to increase in inter electronic repulsion further increasing the screening effect. So the effective nuclear charge reduces resulting in an increase in distance between electron and nucleus.

Eg: Flourine (Z = 9)

	F	F-
Proton	9	9
Electron	9	10
Z/e	9/9 = 1	9/10 = 0.9

As Zeff of fluorine anion (F-) is less than flourine atom (F)

Size of fluorine anion is greater than the parent atom i.e.

F- > F

Due to above mentioned reasons, cations are smaller and anions are larger in size than parent atom.

Atomic radius of chlorine is 175pm.

Read more :

Ionic size or Size of isoelectronic species:-

Chemical species containing same number of electrons but having different nuclear charge are called isoelectronic species. Atomic radius is inversely proportional to effective nuclear charge and hence, as the atomic radius increases, effective nuclear charge decreases for isoelectronic species.

Species	K+	Ca+2	S-2	Cl-
Z	19	20	16	17
e	18	18	18	18
Z/e	19/18	20/18	16/18	17/18

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Periodic trends in ionic radius-

Periodic trends in ionic radius down the group:-

Ionic radius increases down the group in a periodic table. This happens because of the addition of extra shell ( number of electrons) by atoms.

Ions	Configuration	Ionic Radii (nm)
Li+	2	0.076
Na+	2,8	0.102
K+	2,8,8	0.138

Ions	Configuration	Ionic Radii (nm)
F-	2,8	0.133
Cl-	2,8,8	0.181
Br-	2,8,18,8	0.196

Periodic trends in ionic radius across the period:-

Let us consider an example for better understanding of tis concept-

Let us consider 3^rd period. We observe that atomic radius tends to increase first and then decreases gradually and then again increase and slowly decreases down. This happens because the initial elements have a tendency of forming cations and the elements present towards the end of the periodic table tend to form anions.

Period 3	Na+	Mg+	Al+3	P-3	S-2	Cl-
No. of protons	11	12	13	14	15	16
Electronic Configuration	2,8	2,8	2,8	2,8,8	2,8,8	2,8,8
Ionic radius	0.102	0.072	0.054	0.212	0.184	0.181

Factors Affecting Atomic Size:-

In a Period-

Atomic radius ∝ 1/Zeff ∝ Negative Charge/Positive

Charge

Li<Na<K<Rb<Cs

In a Group-

Atomic radius ∝ number of shells

Atomic size decreases from left to right in a period as Zeff increases.

LI > Be > B > C > N > O > F

Atomic size decreases from top to bottom in a group as number of shells are increased.

Li < Na < K < Rb < Cs

Lanthanide Contraction-

Outermost electronic configuration of inner transition elements is (n-2)f1-14,(n-1)s2p6d0-1,ns2 (n=6 or 7)
Electrons enters in (n-2)f orbitals.
Because of complicated structure of f orbital and due to poor shielding f electrons, the outermost shell electrons get attracted towards the nucleus.
In 1st,2nd,3rd, transition series, radii:- 3d>4d≈5d

size%20decrease

Also read -

S-Block variation-

increase%20size

Also check-

Frequently Asked Questions (FAQs)

1. How do you compare ionic radius to atomic radius?

A cation having lost its electron is smaller in size than Parent atom.

An anion having gained electron is larger in size than

parent atom.

2. What is periodicity?

The regular gradation in properties from top to

bottom and from left to right in a period is called

periodicity.

3. What is the magnitude of ionic radius of of magnesium ?

The ionic radius of magnesium ion (Mg²⁺) is 173 pm.

4. What is ionic radius of Chlorine ion ?

The ionic radius of Chlorine anion (Cl⁻ ) is 175 pm.

5. Which element shows highest ionic radius value?

Francium (Fr) is the element showing highest ionic radius having value 348 pm.

6. How does the atomic radius of elements in group 18 (noble gases) change as you move down the group?

The atomic radius of noble gases increases as you move down group 18. This follows the general trend for groups in the periodic table. As you move down the group, each new element adds a new electron shell, which is further from the nucleus. Although the nuclear charge also increases, the effect of the new shell outweighs this increase, resulting in a larger atomic radius.

7. What is the trend in atomic radii for anions compared to their neutral atoms?

Anions have larger atomic radii than their corresponding neutral atoms. When an atom gains one or more electrons to form an anion, the additional electrons increase electron-electron repulsion. This repulsion, combined with the fact that the nuclear charge remains the same, causes the electron cloud to expand, resulting in a larger atomic radius.

8. How does the atomic radius of metalloids compare to that of metals and non-metals in the same period?

Metalloids generally have atomic radii intermediate between those of metals and non-metals in the same period. As you move from left to right across a period, the atomic radius decreases. Metalloids, being located between metals and non-metals, have atomic radii smaller than most metals but larger than most non-metals in the same period.

9. How does the atomic radius of alkali metals compare to that of halogens in the same period?

Alkali metals have significantly larger atomic radii compared to halogens in the same period. This is due to the general trend of decreasing atomic radius across a period. Alkali metals, being on the far left of the periodic table, have the largest atomic radii in their respective periods, while halogens, near the right side, have much smaller atomic radii.

10. What is the effect of a filled subshell on atomic radius?

A filled subshell generally leads to a smaller atomic radius compared to a partially filled subshell. This is because a completely filled subshell is more stable and symmetrical, allowing for more effective shielding of the nuclear charge. This results in a stronger pull on the outer electrons, leading to a more compact atom.

11. How does the atomic radius of transition metals change across a period?

The atomic radius of transition metals generally decreases slightly across a period, but the change is less pronounced compared to main group elements. This is because transition metals add electrons to inner d-orbitals rather than outer orbitals. The poor shielding effect of d-electrons results in a gradual increase in effective nuclear charge and a slight decrease in atomic radius across the transition series.

12. How does the atomic radius of post-transition metals compare to that of transition metals in the same period?

Post-transition metals generally have larger atomic radii compared to transition metals in the same period. This is because post-transition metals have filled d-subshells, which provide better shielding of the nuclear charge compared to the partially filled d-subshells of transition metals. The increased shielding results in a lower effective nuclear charge and, consequently, a larger atomic radius.

13. How does the atomic radius of a cation compare to its neutral atom?

A cation has a smaller atomic radius than its neutral atom. When an atom loses one or more electrons to form a cation, it has fewer electrons but the same number of protons. This results in a higher effective nuclear charge, pulling the remaining electrons closer to the nucleus and reducing the overall atomic radius.

14. How does electron shielding affect atomic radius?

Electron shielding occurs when inner electron shells partially block the nuclear charge experienced by outer electrons. This shielding effect reduces the effective nuclear charge on the outermost electrons, allowing them to be held less tightly to the nucleus. As a result, increased electron shielding tends to increase the atomic radius.

15. How does the atomic radius of a diatomic element compare to its monatomic form?

The atomic radius of a diatomic element is typically smaller than its monatomic form. In a diatomic molecule, the atoms share electrons, creating a bond that pulls the nuclei closer together. This results in a smaller effective radius for each atom in the molecule compared to the isolated atom.

16. How do atomic radii of isotopes compare?

Isotopes of an element have the same number of protons but different numbers of neutrons. The atomic radius of isotopes is generally very similar because the electron configuration, which primarily determines the atomic radius, remains the same. Any slight differences are due to the small effect of nuclear size on electron distribution.

17. Why do elements in the same period have different atomic radii despite having the same number of electron shells?

Elements in the same period have different atomic radii because of the increasing nuclear charge as you move from left to right. While they have the same number of electron shells, the higher number of protons in elements on the right side of the period exerts a stronger pull on the electrons, resulting in a smaller atomic radius.

18. How does the concept of electron affinity relate to atomic radius?

Electron affinity and atomic radius are generally inversely related. Elements with smaller atomic radii tend to have higher electron affinities. This is because smaller atoms can more effectively attract an additional electron due to the closer proximity of the nucleus and the existing electrons.

19. How does the atomic radius change as you move down a group in the periodic table?

As you move down a group in the periodic table, the atomic radius generally increases. This is due to the addition of new electron shells as the atomic number increases. Each new shell is further from the nucleus, leading to a larger overall atomic radius.

20. What is the relationship between atomic number and atomic radius within a group?

Within a group, as the atomic number increases, the atomic radius generally increases. This is because each element in the group has an additional electron shell, which is located further from the nucleus. The effect of increased nuclear charge is outweighed by the addition of new shells, resulting in a larger atomic radius.

21. Why do elements in the same group have similar chemical properties despite having different atomic radii?

Elements in the same group have similar chemical properties because they have the same number of valence electrons in their outermost shell. These valence electrons primarily determine an element's chemical behavior. While atomic radii increase down a group, the valence electron configuration remains similar, leading to comparable chemical properties.

22. What is the relationship between atomic radius and metallic character?

Atomic radius and metallic character are generally directly related. Elements with larger atomic radii tend to exhibit stronger metallic character. This is because larger atoms typically have lower ionization energies and are more likely to lose electrons, a characteristic of metallic behavior. As you move down a group, both atomic radius and metallic character tend to increase.

23. What is the relationship between atomic radius and ionization energy?

Atomic radius and ionization energy are inversely related. As the atomic radius decreases, the ionization energy generally increases. This is because smaller atoms hold their electrons more tightly due to the closer proximity to the nucleus, requiring more energy to remove an electron.

24. What is the periodic trend for atomic radius across a period in the modern periodic table?

Across a period from left to right, the atomic radius generally decreases. This is because as you move across a period, the number of protons in the nucleus increases, leading to a stronger nuclear charge. This increased nuclear charge pulls the electrons closer to the nucleus, resulting in a smaller atomic radius.

25. Why do atoms of noble gases have smaller atomic radii compared to the atoms of alkali metals in the same period?

Noble gases have smaller atomic radii compared to alkali metals in the same period because noble gases have a full outer electron shell. This complete shell experiences a stronger effective nuclear charge, pulling the electrons closer to the nucleus. Alkali metals, on the other hand, have only one electron in their outer shell, which is less tightly bound to the nucleus.

26. Why do transition metals have smaller atomic radii compared to main group elements in the same period?

Transition metals have smaller atomic radii compared to main group elements in the same period because they add electrons to inner d-orbitals rather than outer s- or p-orbitals. The d-orbitals are less effective at shielding the outer electrons from the nuclear charge, resulting in a stronger pull on the electrons and a smaller atomic radius.

27. Why is there a significant decrease in atomic radius between nitrogen and oxygen in the periodic table?

The significant decrease in atomic radius between nitrogen and oxygen is due to the increase in effective nuclear charge. Oxygen has one more proton than nitrogen, but the added electron goes into the same shell. The increased nuclear charge pulls the electrons closer to the nucleus, overcoming the slight increase in electron repulsion and resulting in a smaller atomic radius.

28. How does the concept of effective nuclear charge relate to atomic radius?

Effective nuclear charge is the net positive charge experienced by an electron in an atom, taking into account the shielding effect of inner electrons. As the effective nuclear charge increases, it pulls the electrons closer to the nucleus, resulting in a smaller atomic radius. Conversely, a decrease in effective nuclear charge leads to a larger atomic radius.

29. How does the concept of isoelectronic species relate to atomic radii?

Isoelectronic species are atoms or ions with the same number of electrons. Among isoelectronic species, the atomic or ionic radius decreases as the nuclear charge increases. This is because a higher nuclear charge exerts a stronger pull on the same number of electrons, resulting in a more compact electron distribution and a smaller radius.

30. What is the lanthanide contraction, and how does it affect atomic radii?

The lanthanide contraction is the decrease in atomic and ionic radii across the lanthanide series (elements 57-71). This contraction occurs because the added electrons go into the 4f subshell, which poorly shields the outer electrons from the increasing nuclear charge. As a result, the atomic radii of elements following the lanthanides are smaller than expected.

31. Why do elements in the f-block (lanthanides and actinides) have similar atomic radii despite increasing atomic numbers?

Elements in the f-block (lanthanides and actinides) have similar atomic radii because the added electrons go into the inner f-orbitals. These f-electrons are poor at shielding the outer electrons from the increasing nuclear charge. As a result, the increase in nuclear charge is largely balanced by the poor shielding, leading to only small changes in atomic radius across the f-block series.

32. What is the effect of relativistic effects on the atomic radii of heavy elements?

Relativistic effects become significant for heavy elements, particularly those in the sixth and seventh periods. These effects cause the innermost electrons to move at speeds approaching the speed of light, increasing their mass and causing them to orbit closer to the nucleus. This contraction of inner orbitals leads to better shielding of the nuclear charge, resulting in larger-than-expected atomic radii for the outermost electrons of heavy elements.

33. What is the relationship between atomic radius and electronegativity?

Atomic radius and electronegativity are generally inversely related. As the atomic radius decreases, electronegativity tends to increase. This is because smaller atoms hold their electrons more tightly and have a greater ability to attract electrons in a chemical bond. Elements with small atomic radii, such as fluorine and oxygen, tend to have high electronegativity values.

34. How does the concept of atomic radius relate to the formation of chemical bonds?

Atomic radius plays a crucial role in chemical bond formation. When atoms form covalent bonds, their atomic radii overlap, and the distance between nuclei is typically less than the sum of their individual atomic radii. In ionic bonds, the size difference between cations and anions influences the strength of the electrostatic attraction. Understanding atomic radii helps predict bond lengths, strengths, and the overall geometry of molecules.

35. Why is there a slight increase in atomic radius from argon to potassium, despite potassium being in the next period?

The slight increase in atomic radius from argon to potassium occurs because potassium starts a new period and a new principal energy level. While argon has a complete outer shell (3p6), potassium adds its single valence electron to a new 4s orbital. This new orbital is significantly larger than the previous shell, outweighing the effect of increased nuclear charge and resulting in a larger atomic radius.

36. How does the atomic radius of elements in the s-block compare to those in the p-block of the same period?

Elements in the s-block generally have larger atomic radii compared to elements in the p-block of the same period. This is because s-block elements have their outermost electrons in spherical s-orbitals, which are more diffuse and extend further from the nucleus. P-block elements, with electrons in p-orbitals, experience a stronger effective nuclear charge, resulting in smaller atomic radii.

37. Why is there a more significant change in atomic radius between groups 2 and 13 compared to other adjacent groups?

The significant change in atomic radius between groups 2 and 13 is due to the start of a new block in the periodic table. Group 2 elements have their outermost electrons in s-orbitals, while group 13 elements begin filling p-orbitals. The p-orbitals are less penetrating and experience more shielding, resulting in a larger atomic radius for group 13 elements compared to group 2 elements in the same period.

38. What is the relationship between atomic radius and the strength of metallic bonds?

There is generally an inverse relationship between atomic radius and the strength of metallic bonds. As the atomic radius decreases, the strength of metallic bonds tends to increase. This is because smaller atoms allow for closer packing of nuclei and a higher electron density in the "sea of electrons" characteristic of metallic bonding. This results in stronger electrostatic attractions and, consequently, stronger metallic bonds.

39. How does the atomic radius of noble gases compare to that of other elements in the same period?

Noble gases generally have the smallest atomic radii among the elements in their respective periods. This is because noble gases have completely filled outer electron shells, which experience the maximum effective nuclear charge for that period. The strong attraction between the nucleus and the electrons in the filled shell results in a more compact atom compared to other elements in the same period.

40. How does the atomic radius of elements in the third period compare to those in the second period?

Elements in the third period have larger atomic radii compared to their counterparts in the second period. This is because third-period elements have an additional electron shell (n=3) compared to second-period elements (n=2). The presence of this extra shell results in a greater distance between the nucleus and the outermost electrons, leading to larger atomic radii.

41. How does the atomic radius of diamagnetic elements compare to that of paramagnetic elements in the same period?

Diamagnetic elements generally have smaller atomic radii compared to paramagnetic elements in the same period. This is because diamagnetic elements often have completely filled subshells, which result in more effective shielding and a stronger effective nuclear charge. Paramagnetic elements, with their unpaired electrons, may have partially filled subshells that provide less effective shielding, leading to slightly larger atomic radii.

42. Why do some elements deviate from the general trend of decreasing atomic radius across a period?

Some elements may deviate from the general trend of decreasing atomic radius across a period due to electron configuration effects. For example, the atomic radius of oxygen is slightly larger than that of nitrogen, despite the general trend. This is because oxygen has four electrons in its 2p orbital, leading to increased electron-electron repulsion compared to nitrogen's three 2p electrons, which slightly expands the atomic radius.

43. What is the relationship between atomic radius and melting point for elements in the same group?

For elements in the same group, there is generally an inverse relationship between atomic radius and melting point. As you move down a group, the atomic radius increases, but the melting point often decreases. This is because larger atoms typically have weaker interatomic forces due to the increased distance between nuclei, resulting in lower melting points.

44. Why is there a larger difference in atomic radii between periods 3 and 4 compared to periods 4 and 5?

The larger difference in atomic radii between periods 3 and 4 compared to periods 4 and 5 is due to the introduction of the d-block elements. Between periods 3 and 4, a new principal quantum number is introduced along with the d-orbitals, resulting in a significant increase in atomic size. Between periods 4 and 5, the d-orbitals are already present, so the increase in atomic radius is less pronounced.

45. How does the atomic radius of an element in its ground state compare to its excited state?

An atom in an excited state generally has a larger atomic radius compared to its ground state. In an excited state, one or more electrons occupy higher energy orbitals, which are typically further from the nucleus. This increased distance from the nucleus results in a larger overall atomic radius for the excited atom compared to its ground state configuration.

46. How does the atomic radius of elements in group 14 (carbon group) change as you move down the group?

The atomic radius of elements in group 14 (carbon group) generally increases as you move down the group. This follows the overall trend for groups in the periodic table. Each successive element adds a new electron shell, which is further from the nucleus. Although the nuclear charge also increases, the effect of the new shell outwe